Olympic High School  AP Chemistry
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Period ___ Date ___/___/___
20  Electrochemistry
OXIDATION-REDUCTION EQUATIONS
The Half-Reaction Method:
1.
Write the equation as two half-reactions. Include the particles (atoms, ions, molecules) that are involved in
change of oxidation state.
2.
Balance each half-reaction with respect to atoms and charges; first atoms other than H and O, then O with
H2O and H with H+, and ionic charges with electrons (e).
3.
Equalize the number of electrons lost in the oxidation half-reaction with the number of electrons gained in the
reduction half-reaction.
4.
Add the two half-reactions to form a balanced net ionic equation.
5.
(Basic solution) Add OH ions to each side of the equation to neutralize H+ ions. Cancel H2O molecules.
Example 1:
HCl + K2Cr2O7  KCl + CrCl3 + H2O + Cl2
Example 2:
FeCl2 + KMnO4 + HCl  FeCl3 + KCl + MnCl2 + H2O
Example 3:
S2 + MnO4  S + MnO2 (basic solution)
Example 4:
CuS + NO3  Cu2+ + S + NO (acidic)
PROBLEMS:
1.
2.
Balance these equations:
HNO3 + S  NO2 + H2SO4 + H2O
KMnO4 + HCl + H2S  KCl + MnCl2 + S + H2O
3.
FeCl3 + H2S  FeCl2 + HCl + S
4.
Cu + HNO3  Cu(NO3)2 + NO2 + H2O
5.
NaCl + H2SO4 + MnO2  Na2SO4 + MnSO4 + H2O + Cl2
6.
HMnO4 + HCl  MnCl2 + H2O + Cl2
7.
MnO4 + SO32  Mn2+ + SO42
(acidic)
8.
H2O2 + I  H2O + I2
(acidic)
9.
AsO33 + I2  AsO43 + I
(basic)
10.
Cr + ClO4  CrO2 + ClO3
(basic)